2 definition

• Ionic bond

  Strong electrostatic attraction between two oppositely charged ions. Strength of attraction depends on the relative sizes and charges of ions.
• Cation


  A positively charged ion
• Anion


  A negatively charged ion
• Isoelectronic species

  Chemical species that have the same number of electrons
• Isoelectronic species

  • N3-, O2-, F- ions (10 e-'s)
  • CO and N2 molecules (14 e-'s)
• Covalent bond

  The strong electrostatic attraction between two nuclei and the shared pair of electrons between them
• Polar covalent bond

  Occurs when there is an asymmetric electron distribution within the covalent bond due to difference in electronegativities
• Sigma (σ) bond

  A bond that results from a direct end-on overlap of two orbitals
• Pi (π) bond

  A bond that is formed when two orbitals overlap sideways
• Dative covalent bonding

  Occurs when one atom donates both electrons in a bond
• Shapes of molecules

  Shapes adopted by the molecules so as to minimise the electronic repulsion
• Allotropes

  Different forms of the same element
• Malleable

  A malleable substance can be shaped
• Ductile

  A ductile substance can be drawn into a wire
• Intermolecular forces

  Forces between the molecules (cf. bonding, an intramolecular force)
• Electronegativity

  The ability of an atom to attract the bonding electrons in a covalent bond. The most electronegative elements are small and have a relatively high nuclear charge.
• Dipole

  Difference in charge between the two atoms of a covalent bond caused by a shift in electron density in the bond due to the electronegativity difference between elements participating in bonding
• Polar molecules

  Exist as dipoles
• Metallic bonding

  Strong electrostatic attraction between metal ions and the sea of delocalised electrons that surround them
• Delocalised electrons

  The electrons that are not contained within a single atom or a covalent bond
• Bond length

  Internuclear distance between two covalently bonded atoms
• London forces

  Weak intermolecular forces arising due to fluctuations of electron density within a nonpolar molecule. These fluctuations may temporarily cause the asymmetric electron distribution: the molecule becomes an instantaneous dipole. This dipole can induce a dipole in another molecule, and so on. The attraction increases with size/shape (points of contact between the molecules) and number of electrons (more fluctuations = more instantaneous/induced dipoles).
• Permanent dipole-dipole interactions

  Dipole-dipole attractions between polar molecules. Stronger than London forces.
• Hydrogen bond

  A type of intermolecular force (with some bonding character) between a hydrogen bonded to a more electronegative atom than hydrogen (usually N,O,F) and other atom in a same/different molecule. Directional nature - the bond angle is often 180°. Responsible for anomalous properties of water, e.g. the density of ice < density of water. Ice occupies greater volume than water due to the directional nature of hydrogen bonds within the solid structure.